Chem Winter Final

Intensive and Extensive Properties

  • Intensive Properties:

    • Do not depend on the amount of substance.

    • Examples: density, melting point, boiling point, color.

  • Extensive Properties:

    • Vary with the amount of substance.

    • Examples: mass, volume, length.

Phase Changes

  • Transformations between states of matter:

    • Solid to liquid: melting

    • Liquid to solid: freezing

    • Liquid to gas: evaporation

    • Gas to liquid: condensationi

    • Solid to gas: sublimation

    • Gas to solid: deposition

Physical Changes vs Chemical Changes

  • Physical Change: Change in appearance but not in composition.

  • Chemical Change: Substance transforms into another substance.

Mixtures and Pure Substances

  • Miixture: Varies in composition; may be heterogeneous or homogeneous.

    • Heterogeneous Mixture: Distinct parts with different properties.

    • Homogeneous Mixture: Same composition throughout.

  • Pure Substances: Have a fixed composition.

Democritus

  • Proposed the universe is composed of invisible units called atoms (400 BC).

Dalton's Atomic Theory (John Dalton)

  1. All elements consist of indivisible particles called atoms.

  2. Atoms of the same element are identical; different from those of other elements.

  3. Atoms of different elements combine in simple whole-number ratios for compounds.

  4. In chemical reactions, atoms are rearranged but not changed into different elements.

  • "Billiard Ball" Model (1804): Matter made of small, spherical particles, later proven to be divisible.

Thomson's Model

  • "Plum Pudding" or "Chocolate Chip Cookie" Model:

    • Proposed embedded charges in matter.

    • Discovered the electron in 1897 using a cathode ray tube.

    • Positive charge ("dough") with negative electrons ("chocolate").

Nuclear Model of the Atom

  • Ernest Rutherford discovered flaws in previous atomic models via the Gold Foil Experiment.

  • Gold Foil Experiment:

    • Particles shot through gold foil; most pass straight through, some deflected.

    • Conclusion: Atoms consist mostly of empty space with a small, dense nucleus.

Structure of the Atom

  • Nucleus: Contains protons (positive) and neutrons (neutral).

  • Electrons: Responsible for an atom's chemical properties.

  • Mass and Stability: Neutrons and protons contribute to atomic mass and stability.

  • Identity of Atom: Determined by the number of protons.

Isotopes and Atomic Mass

  • Molecule: Composed of two or more atoms bonded together.

  • Isotopes: Same element, different masses (varying neutrons), identical chemical properties. these

  • Atomic Mass: Average of isotopes considering relative abundances.

Ions

  • Definition: Charged particles resulting from the loss or gain of electrons.

  • Cations: Metals lose electrons, becoming positively charged.

  • Anions: Nonmetals gain electrons, becoming negatively charged.

Moles in Chemistry

  • Mole (mol): Represents $6.022 \times 10^{23}$ particles of a substance (Avogadro's number).

Converting Atoms to Moles

  • Conversion Factor: $1 \text{ mol} = 6.022 \times 10^{23}$ atoms.

    • Example Calculation:

    • For $1.25 \times 10^{23}$ atoms of Mg:
      moles = \frac{1.25 \times 10^{23} \text{ atoms Mg}}{6.022 \times 10^{23} \text{ atoms Mg/mol}} = 0.207 \text{ mol Mg}

  • Example #2:

    • For $0.25$ moles of NO2:
      0.25 \text{ mol NO2} \times 6.022 \times 10^{23} \text{ molecules/mol} = 1.5 \times 10^{23} \text{ molecules NO2}

  • Example #3:

    • For $3.84 \times 10^{24}$ molecules of CO:
      3.84 \times 10^{24} \text{ molecules CO} \times \frac{1 \text{ mol CO}}{6.022 \times 10^{23} \text{ molecules CO}} = 6.37 \text{ mol CO}

Converting Moles to Atoms/Molecules

  • Conversion process: moles → molecules → atoms

  • Known: 2.12 mol C3H8

  • Conversion factors:

    • 1 mol C3H8 = $6.022 \times 10^{23}$ molecules

    • 1 molecule C3H8 = 11 atoms (3 C, 8 H)

  • Calculation:

    • $2.12 \text{ mol C3H8} \times \frac{6.022 \times 10^{23} \text{ molecules}}{1 \text{ mol C3H8}} \times \frac{11 \text{ atoms}}{1 \text{ molecule C3H8}} = 1.40 \times 10^{25} \text{ atoms}$

Mass of a Mole

  • Atomic mass in grams = mass of a mole

  • Molar Mass Examples:

    • Carbon (C): 12.01 g/mol

    • Hydrogen (H): 1.01 g/mol

    • Diatomic Hydrogen (H2): 2.02 g/mol

  • Molar mass of Water (H2O): 18.02 g/mol

    • Calculation: H2 = $2 \times 1.01 = 2.02$; O = $1 \times 16 = 16$

Mass to Molecules

  • Example: Find molecules in 5g of H2O

    1. Convert mass to moles:

    • $5\text{g H2O} \times \frac{1 \text{ mol H2O}}{18 \text{ g}} = 0.28 \text{ mol H2O}$

    1. Convert moles to molecules:

    • $0.28 \text{ mol H2O} \times 6.022 \times 10^{23} \text{ molecules} = 1.673 \times 10^{23} \text{ molecules H2O}$

Molecules to Mass

  • Example: Find grams in 1.25 x 10²⁴ molecules of CO2

    1. Convert molecules to moles:

    • $1.25 \times 10^{24} \text{ molecules} \times \frac{1}{6.022 \times 10^{23} \text{ molecules}} = 2.07 \text{ mol CO2}$

    1. Convert moles to mass:

    • $2.07 \text{ mol CO2} \times 44 \text{ g/mol} = 91.33 \text{ g CO2}$

Reaction Types

  • Combination Reaction (Synthesis):

    • Formula: A + B → AB

    • Example: Na + Cl2 → NaCl

  • Decomposition:

    • Formula: AB → A + B

    • Example: H2O → H2 + O2

  • Single Displacement:

    • Formula: C + AB → CB + A

    • Example: Fe + CuSO4 → FeSO4 + Cu

  • Double Displacement (Ionic):

    • Formula: AB + CD → AD + CB

    • Example: Ca(NO3)2 + NaOH → Ca(OH)2 + NaNO3

Properties of Bonds

  • Metallic Bond:

    • Electrical attraction between valence electrons and nuclei

    • Properties: Lustrous, malleable, ductile, conducts electricity

  • Ionic Compounds:

    • High melting/boiling points, non-conductors in solid state, conduct in molten/aqueous states

  • Covalent Bonding:

    • Formed by sharing electrons; represented by dashes

    • Types: Single (-), Double (=), Triple (≡)

  • Network Covalent Bonding:

    • Electrons shared in multiple directions; strong bond

    • Properties: Extremely high melting points, non-conductors, insoluble in water

    • Examples: Diamond, SiO2 (Quartz)

Molecular Covalent Compounds

  • Valence electrons shared among non-metal atoms in separate molecules.

  • Properties:

    • Relatively low melting point.

    • May not easily dissolve in water.

    • Composed of non-metal atoms.

  • Examples: water ($H2O$), Aspirin ($C9H{10}O4$), Ethanol ($C2H6O$), sugar ($C{12}H{22}O{11}$), butane ($C4H{10}$), Paraffin Wax ($C{20}H_{42}$).

Polar Covalent Bond

  • Formed between atoms with differing electronegativities.

  • Electronegativity: ability to attract electrons.

    • Trend: Increases left to right, decreases top to bottom.

  • Example comparison: Cl > H, O > C.

Nonpolar Covalent Bond

  • Formed by equal sharing of electrons between identical nonmetallic atoms (zero electronegativity difference).

  • Examples: $H2$, $F2$, $O_2$ (BrINClHOF).

Bond Type Prediction

  • Use electronegativity difference:

    • Example: C and Cl ($2.5 - 3 = 0.5$, polar).

    • F and F ($4 - 4 = 0$, nonpolar).

    • Mg and O ($1.2 - 3.0 = 1.8$, ionic).

Ionic Bonds

  • Formed by electron transfer for stability.

  • Typically between Groups 1, 2, 3, and 6 or 7.

  • Examples: NaCl, MgCl2.

    • Formation: Na + Cl → Na+ + Cl- → NaCl.

    • Formation: Mg + Cl → Mg2+ + Cl- → MgCl2.

Definition of Intermolecular Forces (IMF)

  • Attractive forces between molecules.

  • Weaker than chemical bonds (Van der Waals forces).

  • Named after Johannes D. Van der Waals, a Dutch physicist.

Types of IMF:

London Dispersion Forces

  • Attraction between 2 instantaneous dipoles.

  • Caused by asymmetrical electron distribution.

  • Occurs in all atoms and molecules.

Dipole-Dipole Forces

  • Attraction between 2 permanent dipoles.

  • Involves polar molecules.

Hydrogen Bonding

  • Attraction between molecules with N-H, O-H, or F-H bonds.

  • Extremely polar bonds; strongest dipole-dipole interaction.

Strength of IMF

  • Weakest: London Dispersion Forces

  • Medium Strength: Dipole-Dipole Forces

  • Strongest: Hydrogen Bonding

  • Relative strength increases with molar mass (more electrons) and closer molecular proximity.

Ion-Dipole Forces

  • Occur between permanently polar molecules and ionic compounds.

  • Stronger than dipole-dipole forces and similar in nature.

  • Facilitate the dissolving of ionic compounds in polar solvents.

  • Example: Na^+ interacts with the negative end of a polar molecule, Cl^- interacts with the positive end.

  • Strength order: Ion-Dipole > Hydrogen Bond > Dipole-Dipole > London Dispersion.

Rules for Significant Figures

  • Zeros between nonzero digits are significant. (e.g., 50.3 = 3 sig figs)

  • Zeros in front of nonzero digits are not significant. (e.g., 0.892 = 3 sig figs)

  • Zeros at the end of a number and to the right of a decimal are significant. (e.g., 2.0 = 2 sig figs)

  • Trailing zeros in a whole number with a decimal point are significant. (e.g., 520. = 3 sig figs; 520 = 2 sig figs)